The rate of a reaction increases with concentration.
Jon Leary 11SA
Rates of Reactions - Coursework
Hypothesis
The rate of a reaction increases with concentration.
Prediction
I predict that when the concentration of the acid is raised, the rate of the reaction will increase. This is because, when the concentration of acid is higher, more acid particles are present in a given volume of the solution, therefore, more acid particles are available to collide with magnesium particles. This consequently increases the chances of successful collisions (those resulting in a reaction) occurring. As the collision theory states, the more collisions that occur in a period of time, the faster the rate of the reaction. This is because the reaction only occurs when acid and magnesium particles collide and successfully join together to form magnesium chloride and hydrogen particles.
In a previous experiment with acid and marble chips, I found that higher concentrations of acid increased the rate of reaction quite significantly.
About the Reaction
The reaction that we are using to prove the above hypothesis is between an acid and a metal. An acid and a metal will produce a salt and hydrogen when reacted together. Since we are trying to measure the rate of a reaction, a reaction that produces a non-toxic gas is the most suitable as it is easy to see when a given quantity of gas has been produced. The reaction is recorded as beginning when the gas is first produced and stopping when it is no longer produced. However, I think that it would be easier to record the amount of time taken for a certain amount of gas to be produced, as it will be both quicker and more practical, as I may find that quite large amounts of gas are produced. The specific reaction that we will use in this investigation is between hydrochloric acid and magnesium. The products of this reaction are magnesium chloride and hydrogen. The equations for the reaction are as follows:
Dilute Hydrochloric Acid + Magnesium Magnesium Chloride + Hydrogen
2HCl(aq) + Mg(s) MgCl2(aq) + H2(g)
Key Variables and Fair Testing
There are three other key variables that can affect the rate of a reaction. These are:
* Temperature - an increase in temperature will lead to an increase in particle movement, therefore increasing the number of collisions. It will also increase the energy they have when they collide with other particles, making them more likely to be able to achieve the activation energy (minimum amount of energy needed to break the bonds of the original compound and initiate a reaction) and be able to react. Both of which increase the rate of reaction.
* Surface area - if the surface area of one or more of the solid reactants is increased, then the rate of reaction will also increase because there are more particles available to collide at any given moment in time.
* Presence of a catalyst - if a catalyst is present during a reaction (e.g. Manganese (IV) Oxide speeds up the decomposition of Hydrogen Peroxide into Water and Oxygen), the activation energy of the reactants will be lowered, therefore more collisions will result in actual reactions and the rate of the overall reaction will be faster.
In order to stop these key variables and other outside factors from unfairly influencing the rate of the reaction, I will take the following steps:
* Temperature must remain the same during the experiment, because if it changes, so will the rate of reaction.
* The surface area and mass of the magnesium strip must remain the same, because an increase in the surface area would lead to a faster reaction and an increase in mass would mean there are more particles to react, so the reaction would last for longer.
* The volume of ...
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In order to stop these key variables and other outside factors from unfairly influencing the rate of the reaction, I will take the following steps:
* Temperature must remain the same during the experiment, because if it changes, so will the rate of reaction.
* The surface area and mass of the magnesium strip must remain the same, because an increase in the surface area would lead to a faster reaction and an increase in mass would mean there are more particles to react, so the reaction would last for longer.
* The volume of the liquid must be kept the same; otherwise I will be investigating the effect of volume on the rate of reaction, as well as concentration. It will also help to make sure that the concentrations are accurate (e.g. 30ml acid and 20ml water or 10ml acid and 40ml water etc.).
* I must always begin recording the starts and ends of the reactions at pre-determined points.
* To ensure accuracy in the results, I will take an average of three results. If any obviously anomalous results occur, I will redo that specific reaction until the results are satisfactorily accurate.
* I will read all volumes in measuring cylinders at eye level to stop distortion.
Safety
The reactants and products of the reaction are not highly dangerous and were chosen partly because of this reason. However, there are some precautions that must be taken before and during the reaction:
* The hydrochloric acid is corrosive, so extra care should be taken and safety goggles worn when handling it.
* The magnesium ribbon is highly flammable and should not be exposed to a naked flame.
* Neither of the products of the reaction are toxic, the only risk being the high flammability of the hydrogen gas.
* The usual laboratory safety rules should be obeyed, such as not running, ties tucked into shirts etc.
Apparatus
Preliminary Method
In order to determine the quantities of the reactants required, I will conduct a preliminary experiment.
I already know from the previous experiment with marble chips and acid that 50cm³ of acid will be a sensible amount to use. However, I have not used the magnesium before in an experiment like this. Consequently, I do not know how much to use.
I will measure lengths of magnesium rather than weigh it, as it comes in ribbons with a constant width and thickness, therefore the surface area and mass will remain almost the same (magnesium is very light, so, especially with the small quantities used in this reaction, variations in mass will not register on the electronic balance). I will use strips of magnesium with lengths 2cm, 3cm, 4cm, 5cm, 6cm and 8cm and follow the method shown below for the main experiment, but only taking two results for each length (the results do not need to be as accurate). I used 2M concentrated hydrochloric acid, as this will be the highest concentration used in the main experiment, so would take the least time.
Preliminary Results
Concentration of Acid
Length of Magnesium (cm)
Time Taken to Produce 20cm³ of Hydrogen (s)
2M
2
Failed to produce 20cm³ of Hydrogen
2.59
2M
3
8.63
7.28
2M
4
5.81
5.91
2M
5
4.62
4.51
2M
6
3.89
3.95
2M
8
3.23
3.22
0.4M
3
407
0.4M
5
81
Preliminary Analysis
From the results of the experiment, I had decided that 3cm of magnesium would be the optimum length of magnesium to use. This was because 2cm was far too slow and only produced enough gas 50% of the time and 6cm and 8cm were both too quick to accurately measure (small mistakes, such as slightly late starting of the stopwatch would have a bigger impact on the results). I then decided to test how long it would take to produce 20cm³ of hydrogen with 0.4M hydrochloric acid, as this is the lowest concentration of acid molecules that will be used in the main experiment, so would take the most time. I found that it took 407s, a time which is highly impractical when three or more must be done in a lesson and similar time taken for slightly higher concentrations. I then tried 5cm of magnesium at this concentration and found that it produced 20cm³ of hydrogen in 181s - a much more practical length of time. After the second set of results, I decided to use 5cm of magnesium in the main experiment, as it was not too slow, always produced enough gas and was not to quick to accurately record. I will use 50cm³ of solution, as previously stated.
Method
. Measure out 50ml of 2M hydrochloric acid solution in a measuring cylinder and pipette for accuracy and pour into a small conical flask.
2. Cut a 5cm strip of magnesium ribbon using a ruler and clean with emery cloth to remove any dirt that would reduce the surface area and consequently the rate of reaction (see key variables section).
3. Fill the measuring cylinder with water and place upside-down in the water trough.
4. Insert the delivery tube into the measuring cylinder, so that all gas travelling through the delivery tube flows into the measuring cylinder.
5. Drop magnesium into conical flask.
6. Attach a bung and delivery tube to the top of the conical flask.
7. Start stopwatch immediately when hydrogen starts to flow into the measuring cylinder.
8. Stop stopwatch and record time when 20cm³ hydrogen has collected in the measuring cylinder.
9. Clean apparatus and repeat steps 1-8 twice to get a more accurate average reading. Redo any anomalous results.
0. Clean the apparatus and repeat steps 1-9, with the following concentrations of acid:
* 10cm³ water, 40cm³ acid - 1.6M hydrochloric acid
* 20cm³ water, 30cm³ acid - 1.2M hydrochloric acid
* 30cm³ water, 20cm³ acid - 0.8M hydrochloric acid
* 40cm³ water, 10cm³ acid - 0.4M hydrochloric acid
Results
Concentration of Acid (mol/dm³ or M)
Time Taken to Produce 20cm³ of Hydrogen (s)
Average Time (s)
Rate of Reaction (1/sec)
2
4.00
4.38
4.88
4.420
0.249
.6
6.93
7.03
7.53
8.34*
7.163
0.140
.4
2.53
2.84
3.53
2.967
0.077
0.8
26.59*
26.65*
32.72*
34.25
35.65
36.03
35.310
0.028
0.4
90.04*
32.53
37.53
43.28
78.58*
34.447
0.007
I also observed that the conical flask was warm after the reaction.
Line graphs of the results are attached on a separate piece of paper.
All figures are to 3 decimal places.
* Anomalous results that have not been included in the average.
Analysis
The graph clearly shows that as concentration increases, so does the rate of reaction. This complies with the hypothesis and proves the predictions made earlier. The rate of reaction increases with concentration, because as the concentration is raised, so more acid particles are available to come into contact and collide with the magnesium particles. This will mean that more successful collisions (those resulting in a reaction) are likely to occur. The collision theory states that the more collisions between particles there are in a given time, the faster the rate of reaction. The graph shows a very steep curve, indicating that the relationship between the rate of reaction and the concentration of the acid is not directly proportional - if the concentration of the acid is increased by x, the rate of reaction will not increase by x. This is because the reaction is exothermic, (releases heat into the surrounding particles) so the other acid and magnesium particles gain heat energy. By gaining heat energy, they become more active - more collisions occur and the particles have more energy when they collide, so are more likely to have the activation energy required to initiate a reaction, therefore the reaction occurs at a faster rate. However, the curve is only so steep because temperature (one of the key variables mentioned earlier) has been altered. The fact that the reaction is exothermic means that when the reaction already occurs quicker (i.e. when the concentration of acid is greater), the particles next to the ones releasing heat energy gain a lot of heat energy in a short amount of time and have very little time to lose any to their surroundings. Consequently, they react faster than cooler particles and the rate of reaction quickens quite significantly. However, when the rate of reaction already occurs slower (i.e. when the concentration of acid is lower), the particles also gain a lot of heat energy from the ones next to them that are releasing it, but over a longer period of time, because the reaction is already slower and they have time to release it into their surroundings and cool down. Consequently they react closer to the normal speed that would be expected at a constant temperature. This is the reason why the graph curved so steeply. However, if temperature were a constant, I think the graph would be a straight line, as there is no longer any reason why it should be curved, but keeping the temperature the same for all the particles during an exothermic reaction would be very difficult to do. Other evidence that supports this theory is the fact that the conical flasks were hot when we came to empty them, especially the faster reactions (i.e. when the concentration of acid is greater). The hydrogen seemed to collect quicker in the measuring cylinder at the start of the reaction - supporting the theory that the quick release of heat sped up the reaction most at the start, because after that, it would start to escape into the surroundings and the particles would cool down, causing the rate of reaction to slow with them. This is also the point when the concentration of acid is at its highest, because after this point, acid molecules are used up in the reaction.
Evaluation
Overall, the experiment went fairly well - we obtained relatively accurate results that supported the hypothesis and proved the prediction. However, we did obtain some anomalous results (shown by an asterix in the results table). Some were quite far out, but some were only a little out. The ones that were only a little out could have been due to minor inaccuracies in timing or measuring out the quantities of acid, including the water left in the beaker after washing (this would lower the concentration and the rate of reaction) and a small air bubble getting into the measuring cylinder that collected the hydrogen (which would shorten the recorded time for 20cm³ of hydrogen to be collected). These problems could have been solved by timing using electronic probes or just watching more carefully, more careful measuring of the quantities (maybe using a thinner measuring cylinder, so that measurements are easier to see) and drying of the conical flasks before reuse. The more major inaccuracies could have been caused by different concentrations of the acid before it was mixed with the water (I got my acid from many different bottles) or possibly the fact that the magnesium floated on the surface, exposing only half of its total surface area to the acid. These problems could be rectified by only getting acid from one bottle and making sure the magnesium either sinks or floats on the surface for all the reactions to ensure continuity (otherwise, the surface area variable will have been altered - sinking would expose more particles to the acid and floating less). To reduce the effect of the exothermic reaction on the reaction as a whole, the conical flask could be placed in a water bath to help maintain a constant temperature. I could also investigate the effect of temperature separately using different temperature water baths and ice. There is currently no known catalyst for the reaction between magnesium and hydrochloric acid, but the magnesium could be cut up into smaller pieces to increase its surface area. I had to adapt the method after the first few tries, because I had been putting the magnesium in the conical flask before the acid and consequently, hydrogen began to be produced before the bung could be secured. I changed the method to put the acid in first, then just drop the magnesium into it as this is quicker and allows the bung to be secured before any hydrogen can escape.
